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EDEXCEL Chemistry a level > CHEMICAL BONDING AND STRUCTURE > Flashcards

CHEMICAL BONDING AND STRUCTURE Flashcards

Giant structures Discrete molecules Physical properties

1
Q

metallic bonding

A

strong electrostatic attraction between the nuclei of metal cations and delocalised electrons.

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2
Q

metals properties

A 
high mp/bp
high thermal/electrical conductivity
malleable
sonorous
ductile
low IE's
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3
Q

Why do metals have high mp’s?

A

Need to break strong forces of attraction between the cation nuclei and delocalised electrons.
also giant lattice structure requires much energy to break

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4
Q

What determines bp/mp’s?

A

no. of delocalised electrons (group 1 is low while d-block metals are higher)
size of cation (smaller radii has closer electrons to nuclei higher mp)

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5
Q

electrical conductivity in metals

A

movement of delocalised electrons towards the positive terminal of a cell with a potential difference.

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6
Q

thermal conductivity

A

delocalised electrons passing KE along metal.

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7
Q

malleability

A

shape configuration

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8
Q

ductility

A

ability to draw metal into a wire

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9
Q

what does ductility and malleability depend on?

A

the cation and electron movement ability. layers slide over one another and delocalised electrons prevent strong forces of repulsion between cations.

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10
Q

ionic bonding

A

strong electrostatic attraction between oppositely charged ions.

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11
Q

arrangement of ionic substances

A

giant ionic lattice in which electrostatic attraction occurs in all directions.

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12
Q

how do you determine the strength of electrostatic attractions in an ionic substance?

A

calculation of energy per mole of solid needed to separate ions to infinity in which ions can no longer interact

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13
Q

how is strength of ionic bonding affected by size of cation?

A

as cation size increases, amount of energy needed to separate them decreases

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14
Q

how is strength of ionic bonding affected by size of anion?

A

increases as anion size increases

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15
Q

how does size of ions affect energy needed to separate ionic substances to infinity?

A

smaller substances require more energy to separate

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16
Q

why must data of ionic radii all come from the same source?

A

because radii is difficult to measure and several methods could conflict in accuracy.

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17
Q

properties of ionic substances

A

high mp/bp
brittle
water-soluble
poor electrical conductivity as a solid

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18
Q

why do ionic substances have high melting points?

A

giant lattice networks of oppositely charged ions have combined large electrostatic forces.

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19
Q

why are ionic substances brittle?

A

as stress causes ionic layers to slide over one another, meaning same charge ions are next to each other, repelling

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20
Q

why are ionic substances water soluble

A

polarity of water molecues allows separation of ions

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21
Q

reason for electrical conductivity when molten

A

previously fixed electrons become mobile, allowing a migration of electrons towards the positive terminal

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22
Q

what sort of current is used in electrolysis

A

direct

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23
Q

covalent bonding

A

2 atoms overlapping atomic orbitals of which contain a singular electron

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24
Q

3 ways covalent bonds interact

A

sigma (2 s orbitals)
sigma (2 P orbitals)
pi (2 P orbitals)

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25
Q

how may pi orbitals form? how does this affect the molecule?

A

sigma bonds may lead to the formation of pi bonds, creating larger electron density above and below the molecule

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26
Q

bond length

A

the distance between nuclei of 2 atoms covalently bonded.

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27
Q

how does bond length affect covalent bond strength

A

shorter length, greater strength

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28
Q

electronegativity

A

the ability of an atom to attract a bonding pair of electrons.

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29
Q

how does electronegativity change over the periodic table?

A

decreases down groups

increases across periods

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30
Q

What are polar molecule electron densities like?

A

they form asymmetrical electron densities, resulting in one atom having a slight positive charge and the other a slight negative charge (delta minus and delta plus)

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31
Q

what are polar bonds represented by?

A

arrows from one atom to the other

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32
Q

2 bonding ideals

A

100% ionic

100% covalent

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33
Q

polar bond

A

type of covalent bond between 2 atoms where bonding electrons are unequally distributed, meaning slight differing charges of each atom.

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34
Q

Discrete (simple molecules)

A

Electrically neutral group of 2 or more atoms held together by chemical bonds.

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35
Q

Dative covalent bond

A

Bond formed when an empty orbital of one atomoverlaps with an orbital containing a lone pair of electrons.

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36
Q

Dative bond representation

A

Arrow

37
Q

Formation of Al2Cl6

A

Lone pair on chlorine atom overlaps with empty orbital on an aluminium atom.
Only occurs in gaseous state.

38
Q

VSEPR

A

Valence shell electron pair repulsion theory

39
Q

electron pair repulsion theory basis

A

The shape of an electron/ion is caused by repulsion between electron pairs (lone and bond) that surround the central atom, arranging themselves to experience minimal repulsion.

40
Q

order of repulsion strength

A

lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion

41
Q

steps to work out shapes of molecules

A

1.) find no. of areas of electron density (bond pairs/lone pairs) via dot and cross.
Then work out shape

42
Q

2 bond pairs and no lone pair shape

angle?

A

linear

180 degrees

43
Q

3 bond pairs and no lone pairs shape and angle

A

trigonal planar

120 degrees

44
Q

4 bond pairs and no lone pairs shape

angles?

A

tetrahedral

109.5 degrees

45
Q

5 bond pairs and no lone pairs shape

angles?

A

trigonal bipyramidal

90 degrees and 120 degrees

46
Q

6 bond pairs and no lone pairs shape

angles?

A

hexagonal

90 degrees and 180 degrees

47
Q

3 bond pairs and one lone pair
shape? angle?
why?

A

trigonal pyramidal
107 degrees
lone pair-lone pair repulsion is greater than bond pair-lone pair repulsion so the angles is slightly less than 109.5 degrees.

48
Q

2 bond pairs and 2 lone pairs
shape?
angles?

A

v-shaped
104.5 degrees
lone pair-lone pair repulsion > lone pair-bond pair repulsion so bond angles is further depressed.

49
Q

dipole

A

a separation of charge caused by the drift of bonded electrons towards the more electronegative element.

50
Q

why does electronegativity decrease down a group?

A

as atomic radii increases, distance between electrons and nuclei increases, reducing force of attraction between atoms.
also more shielding

51
Q

molecule with intermediate bonding character (both covalent and ionic)

A

calcium carbonate

ionic between calcium and carbonate, covalent within carbonate

52
Q

rough size of an atom

A

10 to the -10 m

53
Q

isoelectronic

A

having the same electronic configuration

54
Q

size of negative ions in relation to atoms that they form from
what about positive?

A

negative ions are larger

positive ions are smaller

55
Q

what is ionic bonding complicated by?

A

ionic packing

covalent character

56
Q

dative bonding

A

covalent bond in which both electrons come from the same atom

57
Q

discrete (simple) molecule

A

electrically neutral group of 2 or more atoms held together by weak chemical bonds

58
Q

DIPOLE

A

when 2 charges of equal magnitude but opposite signs are separated by a small distance

59
Q

factors affecting electronegativity

A

nuclear charge
atomic radius
shielding

60
Q

charge effects on electronegativity

A

higher charge, greater electronegativity

61
Q

atomic radius effect on electronegativity

A

smaller atom, closer the pair of bonding atoms are to the nucleus, meaning stronger force of attraction

62
Q

shielding effect on electronegativity

A

more inner shells, more shielding and so less force of attraction to attract 2 bonding pairs

63
Q

electronegativity down a group

A

decreases
bigger radius
more shielding
less force of attraction

64
Q

electronegativity across a period

A 
increases
smaller radius
higher charge
less shielding
more force of attraction
65
Q

diatomic molecule polarities

A

non-polar

as 2 atoms in each molecule are same and so have same eneg w symmetrical densities

66
Q

symbol representing dipole

A

arrow w cross through

67
Q

linear molecules (symmetrical) polarity

A

non-polar as dipoles cancel each other out

68
Q

trigonal planar molecules polarity

A

non-polar as is symmetrical so dipoles cancel each other out

69
Q

tetrahedral polarity

A

if symmetrical, non-polar

if unsymmetrical, polar as dipoles reinforce one another

70
Q

trigonal pyramidal polarity

A

dipoles reinforce one another so is polar

71
Q

v-shaped molecule polarity

A

dipoles reinforce each other so is polar

72
Q

3 types of intermolecular forces

A

London/dispersion forces
permanent dipole interactions
hydrogen bonds

73
Q

formation of London/dispersion forces

A

created by instantaneous dipoles occurring due to electron density fluctuations, inducing other dipoles between other molecules, aligning so that they interact favourably with one another

74
Q

why do London forces continue to attract each other despite movement

A

as fluctuations in electron density are much quicker than the kinetic energy, so continue to attract each other regardless of energy

75
Q

factors affecting London forces

A

attractive force increases w increasing number of electrons in the molecule (fluctuation in density faster)
more points of contact between molecules, greater overall London force

76
Q

permanent dipole interactions

A

molecules possessing permanent dipoles interact and align, causing forces of attraction between molecules

77
Q

why is London force more significant than permanent dipole interactions

A

random movement of molecules mean they have less alignments and so less interactions
sometimes forces repel so contribute little overall

78
Q

where do london forces exist

A

between all types of molecules, regardless of polarity

79
Q

where do hydrogen bonds exist

A

between all atoms containing an OH group

atom bonded must be more eneg than hydrogen

80
Q

example of hydrogen bonding

A

betweenoxygen atom of one molecule and hydrogen of another

81
Q

h bonding

A

h bonded to an electronegative atom

82
Q

reasons for increasing bp with increasing molecular mass

A

more electrons and so more instantaneous and permanent dipoles form
carbon chain length increase results in more points of contact between adjacent molecules

83
Q

branched chain effects

A

lower bp

due to less points of contact so decrease in IMF

84
Q

anomalous properties of water

A

relatively high bp for little electrons

density of ice is less than water

85
Q

2 conditions needing to be met for a substance to dissolve

A

solute particles must be surrounded by solvent particles

forces of attraction between solvent and solute particles must be strong enough to overcome solute and solute forces

86
Q

van der waal forces

A

sum of al the intermolecular interactions between the molecules, including London and permanent dipole-dipole forces.

87
Q

what happens to the ionic radius of a set of isoelectronic ions as the atomic number increases

A

decreases as the electrons are attracted to the nucleus more strongly.

88
Q

giant ionic lattice

A

repeating units in a regular structure, caused by ions attracted in all directions to oppositely charged ions.

EDEXCEL Chemistry a level (12 decks)

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